chapter 12 gaseous chemical equilibrium
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Chapter 12 Gaseous Chemical Equilibrium. The Concept of Equilibrium. Chemical equilibrium occurs when a reaction and its reverse reaction proceed at the same rate. The Concept of Equilibrium. As a system approaches equilibrium, both the forward and reverse reactions are occurring. - PowerPoint PPT PresentationTRANSCRIPT
Chapter 12GaseousChemical
Equilibrium
The Concept of Equilibrium
Chemical equilibrium occurs when a reaction and its reverse reaction proceed at the same rate.
The Concept of Equilibrium• As a system
approaches equilibrium, both the forward and reverse reactions are occurring.
• At equilibrium, the forward and reverse reactions are proceeding at the same rate.
A System at Equilibrium
Once equilibrium is achieved, the amount of each reactant and product remains constant.
Depicting Equilibrium
In a system at equilibrium, both the forward and reverse reactions are being carried out (it’s a dynamic equilibrium); as a result, we write its equation with a double arrow
N2O4 (g) 2 NO2 (g)
The Equilibrium Constant
• Consider the reaction
• The equilibrium expression for this reaction would be
Kc = [C]c[D]d
[A]a[B]b
aA + bB cC + dD
The Equilibrium Constant
Because pressure is proportional to concentration for gases in a closed system, the equilibrium expression can also be written
Kp =(PC)c (PD)d
(PA)a (PB)b
Relationship between Kc and Kp
• From the ideal gas law we know that
• Rearranging it, we get
PV = nRT
P = RTnV
Relationship between Kc and Kp
Plugging this into the expression for Kp for each substance, the relationship between Kc and Kp becomes
Where
Kp = Kc (RT)n
n = (moles of gaseous product) − (moles of gaseous reactant)
Equilibrium Can Be Reached from Either Direction
As you can see, the ratio of [NO2]2 to [N2O4] remains constant at this temperature no matter what the initial concentrations of NO2 and N2O4 are.
Equilibrium Can Be Reached from Either Direction
This is the data from the last two trials from the table on the previous slide.
Equilibrium Can Be Reached from Either Direction
It does not matter whether we start with N2 and H2 or whether we start with NH3. We will have the same proportions of all three substances at equilibrium.
What Does the Value of K Mean?
• If K >> 1, the reaction is product-favored; product predominates at equilibrium.
What Does the Value of K Mean?
• If K >> 1, the reaction is product-favored; product predominates at equilibrium.
• If K << 1, the reaction is reactant-favored; reactant predominates at equilibrium.
Manipulating Equilibrium Constants
The equilibrium constant of a reaction in the reverse reaction is the reciprocal of the equilibrium constant of the forward reaction.
10.212
=
Kc = = 0.212 at 100C[NO2]2
[N2O4]N2O4 (g) 2 NO2 (g)
Kc = = 4.72 at 100C
[N2O4][NO2]2
N2O4 (g)2 NO2 (g)
Manipulating Equilibrium ConstantsThe equilibrium constant of a reaction that has been multiplied by a number is the equilibrium constant raised to a power that is equal to that number.
Kc = = 0.212 at 100C[NO2]2
[N2O4]N2O4 (g) 2 NO2 (g)
Kc = = (0.212)2 at 100C[NO2]4
[N2O4]22 N2O4 (g) 4 NO2 (g)
Manipulating Equilibrium Constants
The equilibrium constant for a net reaction made up of two or more steps is the product of the equilibrium constants for the individual steps.
A + B C K = 1.9 x 10-4
C + D E + A K = 8.5 x 105
B + D E K =
Heterogeneous Equilibrium
The Concentrations of Solids and Liquids Are Essentially Constant
Both can be obtained by dividing the density of the substance by its molar mass—and both of these are constants at constant temperature.
The Concentrations of Solids and Liquids Are Essentially Constant
Therefore, the concentrations of solids and liquids do not appear in the equilibrium expression
Kc = [Pb2+] [Cl−]2
PbCl2 (s) Pb2+ (aq) + 2 Cl−(aq)
As long as some CaCO3 or CaO remain in the system, the amount of CO2 above the solid will remain the same.
CaCO3 (s) CO2 (g) + CaO(s)
What Are the Equilibrium Expressions for These Equilibria?
Equilibrium Calculations
Equilibrium Calculations
A closed system initially containing
1.000 x 10−3 M H2 and 2.000 x 10−3 M I2
At 448C is allowed to reach equilibrium. Analysis of the equilibrium mixture shows that the concentration of HI is 1.87 x 10−3 M. Calculate Kc at 448C for the reaction taking place, which is
H2 (g) + I2 (g) 2 HI (g)
What Do We Know?
[H2], M [I2], M [HI], M
Initially 1.000 x 10-3 2.000 x 10-3 0
Change
At equilibrium
1.87 x 10-3
[HI] Increases by 1.87 x 10-3 M
[H2], M [I2], M [HI], M
Initially 1.000 x 10-3 2.000 x 10-3 0
Change +1.87 x 10-3
At equilibrium
1.87 x 10-3
Stoichiometry tells us [H2] and [I2]decrease by half as much
[H2], M [I2], M [HI], M
Initially 1.000 x 10-3 2.000 x 10-3 0
Change -9.35 x 10-4 -9.35 x 10-4 +1.87 x 10-3
At equilibrium
1.87 x 10-3
We can now calculate the equilibrium concentrations of all three compounds…
[H2], M [I2], M [HI], M
Initially 1.000 x 10-3 2.000 x 10-3 0
Change -9.35 x 10-4 -9.35 x 10-4 +1.87 x 10-3
At equilibrium
6.5 x 10-5 1.065 x 10-3 1.87 x 10-3
…and, therefore, the equilibrium constant
Kc =[HI]2
[H2] [I2]
= 51
=(1.87 x 10-3)2
(6.5 x 10-5)(1.065 x 10-3)
The Reaction Quotient (Q)
• To calculate Q, one substitutes the initial concentrations of reactants and products into the equilibrium expression.
• Q gives the same ratio the equilibrium expression gives, but for a system that is not at equilibrium.
Only given Initial Concentrations and the value of K
If Q = K,
the system is at equilibrium.
If Q > K,there is too much product and the
equilibrium shifts to the left.
If Q < K,there is too much reactant, and the
equilibrium shifts to the right.
Le Châtelier’s Principle
Le Châtelier’s Principle
“If a system at equilibrium is disturbed by a change in temperature, pressure, or the concentration of one of the components, the system will shift its equilibrium position so as to counteract the effect of the disturbance.”
If an external stress is applied to a system at equilibrium, the system adjusts in such a way that the stress is partially offset as the system reaches a new equilibrium position.
Le Châtelier’s Principle
•Changes in Concentration
N2 (g) + 3H2 (g) 2NH3 (g)
AddNH3
Equilibrium shifts left to offset stress
Le Châtelier’s Principle
•Changes in Concentration continued
Change Shifts the Equilibrium
Increase concentration of product(s) left
Decrease concentration of product(s) right
Decrease concentration of reactant(s)
Increase concentration of reactant(s) right
left
aA + bB cC + dD
AddAddRemove Remove
N2O4(g) 2NO2(g)(colorless) (brown)
Le Châtelier’s Principle
•Changes in Volume and Pressure
A (g) + B (g) C (g)
Change Shifts the Equilibrium
Increase pressure Side with fewest moles of gas
Decrease pressure Side with most moles of gas
Decrease volume
Increase volume Side with most moles of gas
Side with fewest moles of gas
N2O4(g) 2NO2(g)(colorless) (brown)
Is this rxn exo- or endothermic?
Le Châtelier’s Principle
•Changes in Temperature
Change Exothermic Rx
Increase temperature K decreases
Decrease temperature K increases
Endothermic Rx
K increases
K decreases
colder hotter
van’t Hoff equation
• ΔHrxn= enthalpy change for a reaction.
• ΔHrxn = ΣΔHfo
(products) - ΣΔHfo
(reactants) see appendix
ln K1
K2
= ΔHrxn
R [ ]1T2
1T1
-
uncatalyzed catalyzed
Catalyst lowers Ea for both forward and reverse reactions.
Catalyst does not change equilibrium constant or shift equilibrium.
•Adding a Catalyst•does not change K•does not shift the position of an equilibrium system•system will reach equilibrium sooner
Le Châtelier’s Principle
Catalysts increase the rate of both the forward and reverse reactions.
Equilibrium is achieved faster, but the equilibrium composition remains unaltered.
Example
• (e) helium is added and the total pressure increases.
• (f) a catalyst is added
Chemistry In Action
Life at High Altitudes and Hemoglobin Production
Kc = [HbO2]
[Hb][O2]
Hb (aq) + O2 (aq) HbO2 (aq)
Chemistry In Action: The Haber Process
N2 (g) + 3H2 (g) 2NH3 (g) H0 = -92.6 kJ/mol
The Haber Process
The transformation of nitrogen and hydrogen into ammonia (NH3) is of tremendous significance in agriculture, where ammonia-based fertilizers are of utmost importance.
The Haber Process
If H2 is added to the system, N2 will be consumed and the two reagents will form more NH3.
The Haber Process
This apparatus helps push the equilibrium to the right by removing the ammonia (NH3) from the system as a liquid.
Le Châtelier’s Principle
Change Shift EquilibriumChange Equilibrium
Constant
Concentration yes no
Pressure yes (?) no
Volume yes no
Temperature yes yes
Catalyst no no
Adding an inert gas to increase
the total pressure
no no
R(g) + heat B(g) + C(g)
• Increase temperature• Increase [C]• Decrease temperature• Add a catalyst• Decrease [C]• Increase pressure by
pumping in helium• Increase pressure by
decreasing the volume
• Increase blue• Increase red• Increase red• Stays purple• Increase blue• Stays purple
(although diluted a bit)
• Increase red